Modern inorganic chemistry

Transcription

1 Modern inorganic chemistry AN INTERMEDIATE TEXT C. CHAMBERS, B.Sc., Ph.D., A.R.I.C. Senior Chemistry Master, Bolton School A. K. HOLLIDAY, Ph.D., D.Sc., F.R.I.C. Professor of Inorganic Chemistry, The University of Liverpool BUTTERWORTHS

2 THE BUTTERWORTH GROUP ENGLAND Butterworth & Co (Publishers) Ltd London: 88 Kingsway, WC2B 6AB AUSTRALIA Butterworths Pty Ltd Sydney: 586 Pacific Highway, NSW 2U67 Melbourne: 343 Little Collins Street 3000 Brisbane: 240 Queen Street 4000 CANADA Butterworth & Co (Canada) Ltd Toronto: 2265 Midland Avenue, Scarborough, Ontario, M1P 4SL NEW ZEALAND Butterworths of New Zealand Ltd Wellington: Waring Taylor Street 1 SOUTH AFRICA Butterworth & Co (South Africa) (Pty) Ltd Durban: Gale Street First published 1975 Butterworth & Co (Publishers) Ltd 1975 Printed and bound in Great Britain by R..). Acford Ltd., Industrial Estate, Chichester, Sussex.

3 Contents 1 The periodic table 1 2 Structure and bonding 25 3 Energetics 62 4 Acids and bases: oxidation and reduction 84 5 Hydrogen Groups I and II The elements of Group III Group IV Group V Group VI Group VII: the halogens The noble gases The transition elements The elements of Groups IB and IIB The lanthanides and actinides 440 Index 447

4 Preface The welcome changes in GCE Advanced level syllabuses during the last few years have prompted the writing of this new Inorganic Chemistry which is intended to replace the book by Wood and Holliday. This new book, like its predecessor, should also be of value in first-year tertiary level chemistry courses. The new syllabuses have made it possible to go much further in systematising and explaining the facts of inorganic chemistry, and in this book the first four chapters -the periodic table; structure and bonding; energetics: and acids and bases with oxidation and reduction provide the necessary grounding for the later chapters on the main groups, the first transition series and the lanthanides and actinides. Although a similar overall treatment has been adopted in all these later chapters, each particular group or series has been treated distinctively, where appropriate, to emphasise special characteristics or trends. A major difficulty in an inorganic text is to strike a balance between a short readable book and a longer, more detailed text which can be used for reference purposes. In reaching what we hope is a reasonable compromise between these two extremes, we acknowledge that both the historical background and industrial processes have been treated very concisely. We must also say that we have not hesitated to simplify complicated reactions or other phenomena thus, for example, the treatment of amphoterism as a ph-dependent sequence between a simple aquo-cation and a simple hydroxo-anion neglects the presence of more complicated species but enables the phenomena to be adequately understood at this level. We are grateful to the following examination boards for permission to reproduce questions (or parts of questions) set in recent years in Advanced level (A), Special or Scholarship (S), and Nuffield (N) papers: Joint Matriculation Board (JMB). Oxford Local Examinations (O). University of London (L) and Cambridge Local Examina-

5 PREFACE tion Syndicate (C). We also thank the University of Liverpool for permission to use questions from various first-year examination papers. Where appropriate, data in the questions have been converted to SI units, and minor changes of nomenclature have been carried out; we are indebted to the various Examination Boards and to the University of Liverpool for permission for such changes. C.C A.K.H.

6 1 The periodic table DEVELOPMENT OF IDEAS METALS AND NON-METALS We now know of the existence of over one hundred elements. A century ago, more than sixty of these were already known, and naturally attempts were made to relate the properties of all these elements in some way. One obvious method was to classify them as metals and non-metals; but this clearly did not go far enough. Among the metals, for example, sodium and potassium are similar to each other and form similar compounds. Copper and iron are also metals having similar chemical properties but these metals are clearly different from sodium and potassium the latter being soft metals forming mainly colourless compounds, whilst copper and iron are hard metals and form mainly coloured compounds. Among the non-metals, nitrogen and chlorine, for example, are gases, but phosphorus, which resembles nitrogen chemically, is a solid, as is iodine which chemically resembles chlorine. Clearly we have to consider the physical and chemical properties of the elements and their compounds if we are to establish a meaningful classification. ATOMIC WEIGHTS By values of atomic weights (now called relative atomic masses) had been ascertained for many elements, and a knowledge of these enabled Newlands in 1864 to postulate a law of octaves. When the elements were arranged in order ot increasing atomic weight, each

7 2 THE PERIODICTABLE successive eighth element was 4 a kind of repetition of the first'. A few years later, Lothar Meyer and Mendeleef, independently, suggested that the properties of elements are periodic functions of their atomic weights. Lothar Meyer based his suggestion on the physical properties of the elements. He plotted 'atomic volume' the volume (cm 3 ) of the 70 r 60 QJ o < Ll Atomic weight Figure Ll. Atomic volume curve (Lothar Meyer] _j 140 atomic weight (g) of the solid element- against atomic weight. He obtained the graph shown in Figure LL We shall see later that many other physical and chemical properties show periodicity (p. 15). 'VALENCY' AND CHEMICAL PROPERTIES Mendeleef drew up a table of elements considering the chemical properties, notably the valencies, of the elements as exhibited in their oxides and hydrides. A part of Mendeleef s table is shown in Figure 1.2 -note that he divided the elements into vertical columns called groups and into horizontal rows called periods or series. Most of the groups were further divided into sub-groups, for example Groups

8 THE PERIODIC TABLE 3 IA, IB as shown. The element at the top of each group was called the "head' element. Group VIII contained no head element, but was made up of a group of three elements of closely similar properties, called "transitional triads'. Many of these terms, for example group, period and head element, are still used, although in a slightly different way from that of Mendeleef. Group fk A Rb sub- < group Cs r-* vfr* I Li No Cu^i B Ag \ subgroup Ay J HH EZ in ME ITTTf _ Fe Co Ni Ru Rh Pd Os Ir Pt * Francium. unknown to Mendeleef, has been added Figure 1.2. Arrangement oj some elements according to Mendeleef The periodic table of Mendeleef, and the physical periodicity typified by Lothar Meyer's atomic volume curve, were of immense value to the development of chemistry from the mid-nineteenth to early in the present century, despite the fact that the quantity chosen to show periodicity, the atomic weight, was not ideal. Indeed, Mendeleef had to deliberately transpose certain elements from their correct order of atomic weight to make them Hf into what were the obviously correct places in his table; argon and potassium, atomic weights 39.9 and 39.1 respectively, were reversed, as were iodine and tellurium, atomic weights and This rearrangement was later fully justified by the discovery of isotopes. Mendeleef s table gave a means of recognising relationships between the elements but gave no fundamental reasons for these relationships. ATOMIC NUMBER In 1913 the English physicist Moseley examined the spectrum produced when X-rays were directed at a metal target. He found that the frequencies v of the observed lines obeyed the relationship v = a(z ~ b) 2 where a and b are constants. Z was a number, different for each metal, found to depend upon the position of the metal in the periodic table.

9 4 THE PERIODIC TABLE It increased by one unit from one element to the next, for example magnesium 12, aluminium 13. This is clearly seen in Figure 1.3. Z was called the atomic number; it was found to correspond to the charge on the nucleus of the atom (made up essentially of protons and neutrons), a charge equal and opposite to the number of ext ra nuclear Z (atomic number) Figure 1.3. Variation of (frequency]' with Z electrons in the atom. Here then was the fundamental quantity on which the periodic table was built, ATOMIC SPECTRA Studies of atomic spectra confirmed the basic periodic arrangement of elements as set out by Mendeleef and helped to develop this into the modem table shown in the figure in the inside cover of this book. When atoms of an element are excited, for example in an electric discharge or by an electric arc, energy in the form of radiation is emitted. This radiation can be analysed by means of a spectrograph into a series of lines called an atomic spectrum. Part of the spectrum oi hydrogen is shown in Figure 1.4. The lines shown are observed in the visible region and are called the Balmer series after their I/X - figure I A. A part of the atomic spectrum oj hydrogen (/. wavelength)

10 THE PERIODIC TABLE 5 discoverer. Several series of lines are observed, all of which fit the formula where R is a constant (the Rydberg constant). /. the wavelength of the radiation, and n l and n 2 have whole number values dependent upon the series studied, as shown below : Series Lyman Balmer Paschen Brackett , 3, , , , 7, 8 The spectra of the atoms of other elements also consist of similar series, although much overlapping makes them less simple in appearance. THE BOHR MODEL To explain these regularities, the Danish physicist Bohr (again in 1913) suggested that the electrons in an atom existed in certain definite energy levels; electrons moving between these levels emit or absorb energy corresponding to the particular frequencies which appear in the spectrum. As a model for his calculations, Bohr envisaged an atom as having electrons in circular orbits, each orbit corresponding to a particular energy state. The "orbit' model accurately interpreted the spectrum of hydrogen but was less successful for other elements. Hydrogen, the simplest atom, is made up of a proton (nucleus) and an electron. The electron normally exists in the lowest energy state 15 but may be excited from this lowest state, called the ground state, by absorption of energy and reach a higher energy state 2, E 3 always such that the energy change E n is given by E n = const ant / n 2 where n is a whole number called a quantum number. In Bohr's model, the n values corresponded to different orbits, an orbit with radius r l corresponded to n = L r 2 to n = 2 and so on. Improved spectroscopic methods showed that the spectrum of hydrogen contained many more lines than was originally supposed and that some of these lines were split further into yet more lines when

11 6 THE PERIODIC TABLE the excited hydrogen was placed in a magnetic field. An attempt was made to explain these lines using a modified Bohr model with elliptical orbits but this was only partially successful and the model was eventually abandoned. WAVE-MECHANICS With the failure of the Bohr model it was found that the properties of an electron in an atom had to be described in wave-mechanical terms (p. 54). Each Bohr model energy level corresponding to n = 1, 2, 3 is split into a group of subsidiary levels designated by the letters 5, p, d, f. The number n therefore became the number of a quantum level made up of a set of orbitals (p. 54). Interpretation of the effect of a magnetic or electric field on the spectra required that the p, d and / orbitals must also be subdivided so that finally each 'subdivision energy level' can accommodate only two electrons, these being described by the symbols t and j (representing electrons of opposite spin). Each electron can have, therefore, a unique description, its spin and its energy level or orbital. We can summarise the data for the first three quantum levels briefly as shown in Table LI. Table 1.1 ELECTRONS IN THE FIRST THREE QUANTUM LEVELS Orhitnl - - i 2 Quantum level 3 s p d tl tl t! n n Total ti tl Ti Ti n ti n n n Note. The maximum number of electrons that any quantum level can accommodate is seen to be given by the formula 2n 2 where n is the number of the quantum level, for example n 3: the maximum number of electrons is therefore 18. An orbital is characterised by having a single energy level able to accommodate two electrons. The three p orbitals and five d orbitals are given symbols to differentiate them, for example p x, p r p.. representing three orbitals at right angles each capable of containing two electrons.

12 THE MODERN PERIODIC TABLE THE PERIODIC TABLE 7 The close similarity of the atomic spectra of other atoms to that of hydrogen indicates that, as we progressively increase the number of protons in the nucleus and the extranuclear electrons in the atom for a series of elements of increasing atomic number, the additional electrons enter orbitals of the type originally suggested by wavemechanics for hydrogen. The orbitals are filled in order of ascending energy and when several equivalent energy levels are available, each is occupied by a single electron before any pairing of electrons with opposed spin occurs. The order of increasing energy for the orbitals can be deduced from the modern periodic table although for elements of high atomic number (when the electron energy levels are close together) the precise positioning of an electron may be rather uncertain. The filling of the energy levels for the first ten elements, hydrogen to neon, atomic numbers 1-10 is shown in Table 12. Table 1.2 ELECTRONIC CONFIGURATIONS OF THE ELEMENTS HYDROGEN TO NEON Is 2s 2p H He Li Be B C N O F Ne T T I T 1 T T I T 1 T I T 1 T I T I T 1 t! t I T I t 1 T I T 1 T I T T T T T 1 I 1 T T T T T T I T T 4 T I We notice here that the first energy level, quantum number n = 1, is complete at helium and there is only one orbital the Is (first quantum level, s type orbital). When this is full (Is 2 ), we may call it the helium core. Filling of the quantum level begins at lithium; at beryllium the 2s orbital is filled and the next added electron must go into a 2p orbital. All three 2p orbitals have the same energy in the absence of a magnetic or electric field and fill up singly at first elements boron to nitrogen before the electrons k pair up'. (The effect of pairing on the ionisation energy is further discussed on page 16.) The n = 2 quantum level is completed at neon, and again we may use "neon core' for short.

13 8 THE PERIODICTABLE For the next elements, sodium to argon, the n = 3 quantum level fills up in the same way as the n = 2 quantum level. This is shown in Table 1.3. Reference to the modern periodic table (p. (/)) shows that we have now completed the first three periods the so-called ^shorf periods. But we should note that the n = 3 quantum level can still accommodate 10 more electrons. Table 1.3 ELECTRONIC CONFIGURATIONS OF THE ELEMENTS SODIUM TO ARGON Atomic number l.u'ment Is 2s 2p Na n n mm Mg Al Si P S Cl Ar i.e. neon core 3s r n ti Tl n n 3p T Tt TTT T1TT tint mm Notation Ne core 3s 1 Ne core 3s 2 Ne core 3s 2 3p 1 Ne core 3s 2 3p 2 Ne core 3s 2 3/? 3 Ne core 3s 2 3p 4 Ne core 3s 2 3p 5 is 2 2s 2 2p 6 3s 2 3p b The element of atomic number 19 is potassium, strongly resembling both sodium and lithium in its physical and chemical properties. The atomic spectrum of potassium also confirms its position as a Group I element with an electronic configuration resembling that of sodium. These facts indicate that the extra electron in potassium must be placed in a new quantum level and it is therefore ascribed the electronic configuration Ls 2 2.s 2 2p b 3s 2 3p b 4s 1 (i.e. 2, 8, 8, 1). Similar reasoning leads to calcium being given an electronic configuration of Is 2 2s 2 2p 6 3s 2 3p 6 4s 2 (i.e. 2, 8, 8, 2). The following series of 10 elements, atomic numbers inclusive, are all metals, indicating that they probably have the outer electronic configuration of a metal, i.e. 4 or less outer electrons. This is only possible if these electrons are placed in the inner n = 3 quantum level, entering the vacant 3d orbitals and forming a series of transition' metals. We should note that at zinc, atomic number 30, then = 3 quantum level is complete and filling of then = 4 quantum level is resumed with electrons entering the 4p orbitals. The electronic configurations for elements atomic numbers are shown in Table 1.4. Krypton is found to be an extremely unreactive element indicating that it has a stable electronic configuration despite the fact that the n = 4 quantum level can accommodate 24 more electrons in the d and / orbitals.

14 THE PERIODIC TABLE 9 Table 1.4 ELECTRONIC CONFIGURATION OF THE ELEMENTS POTASSIUM TO KRYPTON Atomic Element Is 2s 3s 3p 5d 4s 4p number * * K Ca Sc Ti v Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Argon core T T tl tl n Ti ti Ti Ti Ti tl Tl tl T T n TlnTl t! TI Tl Ti n Ti T f T t T Ti ti TI Ti ti tl n ti t T t t tti tl n nn n t rt T TTi Tl n Tl n nn Ti t Ti Ti Ti n ti n u ti ttin ti n ti n Tl T r T Ti ti Ti T Tt Ti Ti t T T Ti * The tendency to attain either a half filled or fully filled set of d orbitals at the expense of the outer s orbital is shown by both chromium and copper and should be noted. This apparent irregularity will be discussed in more detail in Chapter 13. Note. The electronic configuration of any element can easily be obtained from the periodic table by adding up the numbers of electrons in the various quantum levels. We can express these in several ways, for example electronic configuration of nickel can be written as Is 2 2s 2 2p 6 3s 6 3<i 8 4s 2, or more briefly ('neon core') 3d 8 4s 2, or even more simply as Chemical properties and spectroscopic data support the view that in the elements rubidium to xenon, atomic numbers 37-54, the 5s, 4d 5p levels fill up. This is best seen by reference to the modern periodic table p. (/). Note that at the end of the fifth period the n = 4 quantum level contains 18 electrons but still has a vacant set of 4/ orbitals. The detailed electronic configurations for the elements atomic numbers can be obtained from the periodic table and are shown below in Table 1.5. Note that the filling of the 4/ orbitals begins after lanthanum (57) and the 14 elements cerium to lutetium are called the lanthanides (Chapter 15). The electronic configuration of some of the newly discovered elements with atomic numbers greater than 95 are uncertain as the energy levels are close together. Filling of the 5/ orbitals does begin after actinium (89) and the remaining elements are generally referred to as actinides (Chapter 15).

15 Table 1.5 ELECTRONIC CONFIGURATIONS OF THE ELEMENTS CAESIUM TO LAWRENCIUM llwiii Atomic (wink Is 2s If 3s!f.Id 4s4p4J 4f 5s5pM 5/ fc Cs Ba I I H La Cc Pr Nd Pm Sm h Gd Tb Dy Ho Er Tm Yb LII W I UO HO (2) 26 2 HO (3) 26 2 HO (4) 26 2 HO (5) HO (7) 26 (1) (8) 26 (1) (10) k 10 (11) (12) ( 1 2 (2) 12) (2) (2) (2) (2) (2) I I I Hf Ta W Re Os is f) HO I I I I I

16 H H 9 I 9 I 9 Z 9 I 9 Z 9 Z 9 Z 9 I 9 I 9 Z 9 Z 9 I 9 Z 92 9 Z Z 91 n u n i n 9 z 91 z 91 n i i n i i i Lf> % W 6 16 Dj D 19 toy ^N n 11 9 I S I I I I I I I I 01 n H 01 H W 0! H H 01 H H 01 H H 01 H H 0! 9 I H 0! H H Ml H i U H 9 I 9 Z 9 I 9 Z 9 Z Q 7 Q 7 9 Z 9 I 9 I 9 I 9 Z 9 I 9 Z 9 I 9 Z 9 2 I I I I I I I I I I I I I I I I 08 fit Ji a n»h ny

17 12 THE PERIODICTABLE FEATURES OF THE PERIODIC TABLE 1. Chemical physical and spectroscopic data all suggest a periodic table as shown on p. (/). 2. The maximum number of electrons which a given quantum level can accommodate is given by the formula 2n 2 where n is the quantum level number. 3. Except for the n = 1 quantum level the maximum number of electrons in the outermost quantum level of any period is always eight. At this point the element concerned is one of the noble gases (Chapter 12). 4. Elements in the s and p blocks of the table are referred to as typical elements whilst those in the d block are called "transition elements" and those in the/block are called actinides and lanthanides (or w rare earth' elements). 5. The table contains vertical groups of elements; each member of a group having the same number of electrons in the outermost quantum level. For example, the element immediately before each noble gas, with seven electrons in the outermost quantum level, is always a halogen. The element immediately following a noble gas, with one electron in a new quantum level, is an alkali metal (lithium, sodium, potassium, rubidium, caesium, francium). 6. The periodic table also contains horizontal periods of elements, each period beginning with an element with an outermost electron in a previously empty quantum level and ending with a noble gas. Periods 1, 2 and 3 are called short periods, the remaining are long periods; Periods 4 and 5 containing a series of transition elements whilst 6 and 7 contain both a transition and a 4 rare earth' series. 7. Comparison of the original Mendeleef type of periodic table (Figure 1.2} and the modern periodic table (p. (/)) shows that the original group numbers are retained but Group I, for example, now contains only the alkali metals, i.e. it corresponds to the top two Group I elements of the Mendeleef table together with Group I A. At the other end of the table, Group VII now contains only the halogens, i.e. the original Group VIIB. The transition elements, in which the inner d orbitals are being filled, are removed to the centre of the table and the "rare earth' elements, in which the^/ orbitals are being filled, are placed, for convenience, at the bottom of the table, eliminating the necessity for further horizontal expansion of the whole table. The original lettering of the transition metal groups, for example VIB, VIIB and so on is still used, but is sometimes misleading and clearly incomplete. However, we may usefully refer, for example, to

18 THE PERIODiCTABLE 13 Group IIB and know that this means the group of elements zinc, cadmium and mercury, whilst Group I1A refers to the alkaline earth metals beryllium, magnesium, calcium, barium and strontium. When Mendeleef devised his periodic table the noble gases were unknown. Strictly, their properties indicate that they form a group beyond the halogens. Mendeleef had already used "Group VIIF to describe his "transitional triads' and the noble gases were therefore placed in a new Group O. 8. The transition or d block elements, in which electrons enter inner d orbitals, form a well-defined series with many common and characteristic features. They are all metals; those on the right of the block are softer and have lower melting points than those on the left (Table 13,2, p. 360). Many are sufficiently resistant to oxidation, corrosion and wear to make them useful in everyday life. They have similar ionisation energies (Figure L6\ often give ions of variable valency, and readily form complexes (pp. 46, 362) many of which are coloured. However, regular gradations of behaviour, either across a series or down a group are much less apparent than in the typical s and p block elements. The elements at the end of each transition series copper and zinc in Period 4, silver and cadmium in Period 5 and gold and mercury in Period 6 have d orbitals which are filled. When copper and silver form the copper(i) ion Cu + and the silver ion Ag + respectively, and zinc and cadmium the ions Zn 2+ and Cd 2+ respectively, the inner d orbitals remain filled. Are these elements and ions properly called "transition' elements and ions? We shall see in Chapters 13 and 14 that their properties are in some respects intermediate between those characteristic of a transition metal and a non-transition metal. Thus zinc, for example, is like calcium in some of its compounds but like a transition metal in others. Again, silver has some properties like an alkali metal but also has "transition-like' properties. The elements gold and mercury show little resemblance to any non-transition metals, but their 'transition-like' properties are not much like those of other transition metals either. In the older Mendeleef form of the periodic table, the elements copper, silver and gold often called the 'coinage' metals occupied Group IB, and zinc, cadmium and mercury Group IIB, these being subdivisions of Groups I and II respectively. However, there are no really very good grounds for treating these two trios as groups; copper, silver and gold have few resemblances, and Group IB does not resemble Group IA the alkali metals. These six elements obviously present a problem ; usually they are treated as transition metals or separately as 'the B metals The lanthanides and the subsequently discovered actinides do

19 14 THE PERIODICTABLE not fit into the Mendel eef table and can only be fitted into the modern table by expanding it sideways to an inconvenient degree. They are. therefore, placed separately at the bottom of the table. These two series of elements are now recognised as being inner transition elements, when electrons enter a quantum level two units below that of the outer. Many properties depend upon the outer electronic configurations and hence we can correctly predict that the lanthanides and actinides are two series of closely similar elements. 10. In noting changes of properties down the typical element groups I-VII of the periodic table, it soon becomes apparent that frequently the top or head element in each group does not fall into line with the other elements below it. This is clearly seen when we consider the melting points and boiling points of elements and their compounds (p. 17), and when we come to look at the properties of the individual groups in detail we shall see that the head element and its compounds are often exceptional in both physical and chemical properties. It will be sufficient to note here that all the head elements in Period 2, namely lithium, beryllium, boron, carbon, nitrogen, oxygen and fluorine, have one characteristic in common they cannot expand their electron shells. The elements of Periods 3 onwards have vacant d orbitals, and we shall see that these can be used to increase the valency of the elements concerned but in Period 2 the valency is limited. Unlike 'typical element' groups the 'transition metal' groups do not have head elements. 11. Although the head element of each group is often exceptional in its properties, it does often show a resemblance to the element one place to its right in the period below, i.e. Period 3. Thus lithium resembles magnesium both physically and chemically. Similarly beryllium resembles aluminium and boron resembles silicon but the resemblances of carbon to phosphorus and nitrogen to sulphur are less marked. Oxygen, however, does resemble chlorine in many respects. These are examples of what is sometimes called the diagonal relationship in the periodic table. 12. By reference to the outline periodic table shown on p. (i) we see that the metals and non-metals occupy fairly distinct regions of the table. The metals can be further sub-divided into (a) 'soft' metals, which are easily deformed and commonly used in moulding, for example, aluminium, lead, mercury, (b) the 'engineering' metals, for example iron, manganese and chromium, many of which are transition elements, and (c) the light metals which have low densities and are found in Groups IA and IIA.

20 IMPORTANT PROPERTIES WHICH SHOW A PERIODIC FUNCTION IONISATION ENERGY THE PERIODICTABLE 15 Reference has already been made to Lothar Meyer's plot of "atomic volume' against atomic weight as a demonstration of a physical property of the elements and Figure L5 shows a modem plot of 'atomic volume' against atomic number. Although regularities are clearly observable "atomic volume' has no single meaning for all the elements certainly it does not measure atomic size, a quantity which depends on the state of aggregation of the element. There are, however, more fundamental physical properties which show periodicity. to 60 u o>- 50 4O u I 30 <t 20 IO IO Atomic number Figure 1.5. Atomic volume and atomic number One of these is the first ionisation energy. This is the energy needed to remove one electron from a free atom of the element, i.e. for the process : where M is the element atom. A plot of first ionisation energy against atomic number is shown in Figure 1.6 (units of ionisation energy are kjmor 1 ). Clearly the general tendency is for metals to have low ionisation energies and non-metals to have rather high ionisation energies. We should also note that the first ionisation energies rise as we cross a

21 >2000 o I Hg.Rh Pb Atomic number Figure 1.6. First ionisation energies of the elements CD C O i! /7 th ionisation Figure 1.7. Successive ionisation energies for potassium

22 THE PERIODICTABLE 17 period, although not quite regularly, and fall as we descend a group, for example lithium to caesium. The fall in ionisation energy as we descend a group is associated with the change from non-metallic to metallic character and is very clearly shown by the Group IV elements, carbon, silicon, germanium and tin. Here then is a link between the physico-chemical property ionisation energy and those chemical properties which depend on the degree of metallic (electropositive) character of the elements in the group. If we consider the successive (first, second, third...) ionisation energies for any one atom, further confirmation of the periodicity of the electron quantum levels is obtained. Figure 1.7 shows a graph of Iog 10 (ionisation energy) for the successive removal of 1, 2, 3, electrons from the potassium atom (the log scale is used because the changes in energy are so large). The stabilities of the noble gas configurations at the 18 (argon), 10 )neon) and 2 (helium) levels are clearly seen. The subject of ionisation energies is further discussed in Chapters 2 and 3. MELTING AND BOILING POINTS Both melting and boiling points show some periodicity but observable regularities are largely confined to the groups. In Group O, the noble gases, the melting and boiling points of the elements are low but rise down the group; similarly in Group VIIB, the halogens, the same trend is observed. In contrast the metals of Group IA (and II A) have relatively high melting and boiling points and these decrease down the groups. These values are shown in Figure 1.8. If we look at some of the compounds of these elements we find similar behaviour. Thus the hydrides of Group ynb elements (excepting hydrogen fluoride, p. 52) show an increase in melting and boiling points as we go down the group. These are generally low, in contrast to the melting and boiling points of the Group IA metal chlorides (except lithium chloride) which are high and decrease down the group. The values are shown in Figure 1.9(a) and (b). Clearly the direction of change increase or decrease down the group depends on the kind of bonding. Between the free atoms of the noble gases there are weak forces of attraction which increase with the size of the atom (Chapter 12) and similar forces operate between the molecules of the hydrogen halides HC1, HBr and HI. The forces between the atoms in a metal and the ions in a salt, for example sodium chloride, are very strong and result in high melting and boiling points. These forces decrease with increasing size of atom and ion and hence the fall in melting and boiling points.

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24 19 TOOr Figure 1.8. (a] M.p. and b.p. of Group I A metals, (b) m.p. and b.p. of Group O elements, (c) m.p. and b.p. of the halogens Table 1.6 PERIOD 3 Group I II III IV V VI VII Fluorides Oxides Hydrides NaF Na 2 O NaH MgF 2 MgO MgH, A1F 3 SiF 4, SiO 2 (Am; PF 5 SF 6 (P 2 O 5 ) 2 SO 3 DO CTT i jn ^ on 2 C1F 3 C1 2 0, C1H Group I II in Table 1.7 PERIOD 4 IV V VI VII Fluorides Oxides Hydrides KF K 2 O KH CaF 2 CaO CaH 2 GaF 3 GeF 4 Ga GeO 2 GaH, GeH 4 AsF 5 (As 2 O s ) 2 SeO 3 AsHj ' SeH 2 BrH

25 20 THE PERIODIC TABLE a I H HI LiCl NaCl KCl RbCl CsCl Figure 1.9. (a) M.p. and h.p. of the halogen hydrides HX, (b) m.p. and b.p, of the Group IA chlorides VALENCY Mendeleef based his original table on the valencies of the elements. Listed in Tables L6 and 1.7 are the highest valency fluorides, oxides and hydrides formed by the typical elements in Periods 3 and 4. From the tables it is clear that elements in Groups I-IV can display a valency equal to the group number. In Groups V-VIL however, a group valency equal to the group number (x) can be shown in the oxides and fluorides (except chlorine) but a lower valency (8 x) is displayed in the hydrides. This lower valency (8 x) is also found in compounds of the head elements of Groups V-VIL CHEMICAL CHARACTER In any group of the periodic table we have already noted that the number of electrons in the outermost shell is the same for each element and the ionisation energy falls as the group is descended. This immediately predicts two likely properties of the elements in a group. (a) their general similarity and (b) the trend towards metallic behaviour as the group is descended. We shall see that these predicted properties are borne out when we study the individual groups.

26 THE PERIODIC TABLE 21 Increasing metallic electropositive behaviour down a group also implies a change in the character of the oxides. They will be expected to become more basic as we descend the group and a change from an acidic oxide, i.e. an oxide of a non-metal which readily reacts with OH~ or oxide ions to give oxoacid anions* to a basic oxide, i.e. one which readily yields cations, in some groups. The best example of such a change is shown by the Group IV elements; the oxides of carbon and silicon are acidic, readily forming carbonate and silicate anions, whilst those of tin and lead are basic giving such ions as Sn 2+ and Pb 2+ in acidic solution. Metallic character diminishes across a period and in consequence the oxides become more acidic as we cross a given period. This is clearly demonstrated in Period 3: Na 2 O MgO A1 2 O 3 SiO 2 (P 2 O 5 ) 2 SO 3 C1 2 O 7 + Basic Amphoteric + Acidic > Similar trends are shown by all periods except Period 1. USES OF THE PERIODIC TABLE The most obvious use of the table is that it avoids the necessity for acquiring a detailed knowledge of the individual chemistry of each element. If, for example, we know something of the chemistry of (say) sodium, we can immediately predict the chemistry of the other alkali metals, bearing in mind the trends in properties down the group, and the likelihood that lithium, the head element, may be unusual in certain of its properties. In general, therefore, a knowledge of the properties of the third period elements sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine and argon, is most useful in predicting the properties of the typical elements below Period 3. As regards the transition elements, the first row in particular show some common characteristics which define a substantial part of their chemistry; the elements of the lanthanide and actinide series show an even closer resemblance to each other. One of the early triumphs of the Mendeleef Periodic Table was the prediction of the properties of elements which were then unknown. Fifteen years before the discovery of germanium in 1886, Mendeleef had predicted that the element which he called 'ekasilicon' would be discovered, and he had also correctly predicted many of its properties. In Table 1.8 his predicted properties are compared with the corresponding properties actually found for germanium. Until relatively recently there were other obvious gaps in the

27 22 THE PERIODiCTABLE periodic table, one corresponding to the element of atomic number 87. situated at the foot of Group I A, and another to the element of atomic number 85. at the foot of the halogen group (VIIB). Both of these elements were subsequently found to occur as the products from either natural radioactive decay or from artificial nuclear reactions. Both elements are highly radioactive and even the most stable isotopes have very short half lives; hence only minute quantities of the compounds of either francium or astatine can be accumulated. Table 1.8 PREDICTED PROPERTIES OF GERMANIUM Property Predicted for Ekusilicon* (Es) Found for Germanium Relative atomic mass Density (gcm~ J ) Colour Heat in air Action of acids Preparation Tetrachloride Dirty grey White EsO, Slight EsO 2 4- Na b.p. 373 K, density 1.9 gem" :>,; k Greyish-white White GeO, None by HCl(aq) Ge0 2 + C b.p. 360 K, density 1.89 gem" 3 Taking francium as an example, it was assumed that the minute traces of francium ion Fr + could be separated from other ions in solution by co-precipitation with insoluble caesium chlorate (VII) (perchlorate) because francium lies next to caesium in Group IA. This assumption proved to be correct and francium was separated by this method. Similarly, separation of astatine as the astatide ion At" was achieved by co-precipitation on silver iodide because silver astatide AgAt was also expected to be insoluble. It is an interesting speculation as to how much more difficult the isolation of these two elements might have been if the periodic classification had not provided us with a very good 'preview' of their chemistries. QUESTIONS 1. What do you regard as the important oxidation states of the following elements: (a) chlorine. (b) lead.

28 (c) sulphur, (d) iron? THE PERIODIC TABLE 23 Illustrate, for each valency given, the electronic structure of a compound in which the element displays that valency. Discuss, as far as possible, how far the valencies chosen are in agreement with expectations in the light of the position of these elements in the Periodic Table. (L, S) 2. How, and why, do the following vary along the period sodium to argon: (a) the relative ease of ionisation of the element, (b) the physical nature of the element, (c) the action of water on the hydrides? (C, A) 3. A century ago, Mendeleef used his new periodic table to predict the properties of 'ekasilicon', later identified as germanium. Some of the predicted properties were: metallic character and high m.p. for the element; formation of an oxide MO 2 and of a volatile chloride MC1 4. (a) Explain how these predictions might be justified in terms of modern ideas about structure and valency. (b) Give as many other 'predictions' as you can about the chemistry of germanium, with reasons. (Liverpool B.Sc.,Part I) 4. The following graph shows the variation in atomic radius with increasing atomic number: 25 r E 20 2 b E 10 Cu Br Atomic number 50 60

29 24 THE PERIODIC TABLE (a) What deduction can you make from this graph? (b) Continue the graph to element 60(Nd), and mark on it the approximate positions of the elements (i) Ag (element 47), (ii) I (element 53), (iii) Ba (element 56) (c) Explain briefly (i) the decrease in atomic radius from Li to F, (ii) the increase in atomic radius from F to Br, (iii) the very large atomic radii of the alkali metals, Li to K. (JMB, A) 5. Give the electronic configurations of elements with atomic numbers, 7,11,17,20,26,30 and 36. In each case give the oxidation state (or states) you expect each element to exhibit. 6. Explain the terms, (a) typical element (b) transition element, (c) rare earth element, (d) group, (e) period, (f) diagonal relationship, as applied to the periodic table of elements. In each case give examples to illustrate your answer.

30 Structure and bonding THE NATURE OF THE PROBLEM A very superficial examination of a large number of chemical substances enables us to see that they differ widely in both physical and chemical properties. Any acceptable theory of bonding must explain these differences and enable us to predict the properties of new materials. As a first step towards solving the problem we need to know something of the arrangement of atoms in chemical substances. The structure of a solid can be investigated using a beam of X-rays or neutrons. From the diffraction patterns obtained it is possible to find the arrangement of the particles of which it is composed. Measurement of the amount of heat needed to melt the solid yields information concerning the forces of attraction between these particles, whilst the effect of an electric current and simple chemical tests on the solid may tell if it is a metal or a non-metal. Should the material be a non-conducting solid, we can determine whether it is composed of ions by investigating the effect of an electric current on the molten material. Results of such investigations suggest that there are four limiting kinds of structure and these will be briefly considered. THE METALLIC LATTICE In a pure metal the atoms of the solid are arranged in closely packed layers. There is more than one way of achieving close packing but it 25

31 26 STRUCTURE AND BONDING is generally true to say that each atom is surrounded by as many neighbouring atoms as can be accommodated in the space available. There are no directed forces between the atoms and each atom 'attracts' as many similar atoms as can be accommodated. The ease with which metals conduct electricity indicates that the electrons are only loosely held in this type of structure. THE GIANT MOLECULE LATTICE This is a relatively rare structure, diamond being probably the best known example. Here, the carbon atoms are not close-packed. Each carbon is surrounded tetrahedrally by four other carbon atoms (Figure 2.1). Clearly, each carbon is exerting a tetrahedrally directed Figure 2.1. Structure of diamond force on its neighbours and such directed forces are operative throughout the whole crystal Diamond is found to be a refractory solid, i.e. it has an extremely high melting point, indicating that the bonding forces are extremely strong. Boron nitride (BN) n and silicon carbide (SiC) n (carborundum) are similar types of solid. These solids are non-conducting, indicating that the electrons are less free and more localised than the electrons in a metal which move easily allowing an electric current to flow through the lattice. THE GIANT IONIC LATTICE This is one of the most familiar types of structure in inorganic chemistry. The crystals can usually be melted in the laboratory

32 STRUCTURE AND BONDING 27 although considerable heating is often required. It can be concluded, therefore, that strong forces exist between the particles comprising the crystals, these being usually intermediate in strength between those found in a metal and those found, for example, in diamond. Although the solid crystals do not conduct electricity, the melt does, indicating that the lattice is comprised of charged species, i.e. ions. These ions carry the current and are discharged at the oppositely charged electrode where the products can be identified. X-ray diffraction studies indicate that the ions form a regular lattice, each ion being surrounded by a number of ions of the opposite charge; this number depends on the sizes of the ions concerned and is not dictated by directed forces of attraction*. We can correctly assume the non-directional forces of attraction holding the ions together to be electrostatic in nature. MOLECULAR CRYSTALS This is a very large group comprising mainly crystalline organic materials, but a number of inorganic substances, for example iodine, also come under this heading. These substances melt easily, and may even sublime, indicating the presence of relatively weak forces. They do not conduct electricity in the solid or fused state indicating that the electrons present are localised in strong bonds. These bonds, however, do not permeate the entire structure, as in diamond, and the crystal is comprised of molecules with strong forces between the constituent atoms, but the intermolecular forces are weak. In substances which are liquid or gaseous at ordinary temperature, the forces of attraction between the particles are so weak that thermal vibration is sufficient for them to be broken. These substances can be converted into solids by cooling to reduce the thermal energy. The above classification of structures is made primarily for convenience. In fact, the structures of many compounds cannot be precisely described under any of these classes, which represent limiting, or ideal cases. However, we shall use these classes to examine further the limiting types of bonding found in them. * Many ions can, of course, contain more than one atom (for example NO 3, SOj ) and directed forces hold together the individual atoms within each of these ionic species.

33 28 STRUCTURE AND BONDING THE ELECTRONIC THEORY OF VALENCY After Dalton, in 1807, had put forward the theory that chemical combination consisted of a union between atoms, chemists began their search for the cause and mechanism of the unions. Many ideas were put forward during the following years but, following the discoveries about the structure of the atom, it was realised that the nuclei of atoms were unaffected by chemical combination and that union of atoms must result from interaction between the extranuclear electrons. Kossel and Lewis, working independently in 1916, recognised that the atoms of the different noble gases, with the one exception of helium, each had an outer quantum level containing eight electrons; they therefore suggested that this arrangement must be connected with stability and inactivity, and that reactions occurred between atoms such that each element attained a noble gas configuration. The rearrangement of electrons into stable octets could occur in two ways: (a) by giving or receiving electrons or (b) by sharing electrons. Since 1916 it has been discovered that some noble gases (originally called the inert gases) do form compounds and also there are many reactions known in which elements do not achieve a noble gas configuration. Nevertheless, the theory was a considerable advance towards modem ideas and provides a good basis for discussion. ELECTRON TRANSFER BONDING ELECTROVALENCY The electronic configuration of any element can quickly be deduced from the periodic table. Consider the reaction, for example, between sodium Is 2 2s 2 2p 6 3s 1 (2,8,1) and chlorine Is 2 2s 2 2p 6 3s 2 3p 5 (2.8.7). The theory tells us that combination will occur by electron transfer from the sodium to the chlorine to produce the noble gas configurations 2,8 (Ne) and 2,8,8 (Ar) respectively. Sodium, atomic number 11, becomes the sodium cation Na +, and chlorine the chloride anion Cl~. Electrostatic attraction between these two ions then holds the compound together. This kind of bonding is found in 'giant ionic lattice' compounds and is an example of electrovalency, the bond being said to be ionic. A full discussion of the chemical energetics of such processes will be found in Chapter 3 but at this point it is desirable to consider the energy changes involved in the electron transfer process. The questions to be answered are briefly: 1. What energy changes occur when an element achieves a noble gas configuration?

34 STRUCTURE AND BONDING How does the ease of ion formation change as we cross the periodic table 3. What changes occur as we descend the groups of the table? Consider first the formation of cations by electron loss. Here the important energy quantity is the ionisation energy. As we have seen (p. 15), the first ionisation energy is the energy required to remove an electron from an atom, i.e. the energy for the process M(g)-»M + (g)4- e~ (1 mole) the second, third and fourth ionisation energies being the additional energies required to remove subsequent electrons from the increasingly positively charged ion, the element and the ions formed all being in the gaseous state. Ionisation energies can be obtained from current-voltage plots for gaseous discharges or more conveniently and completely from spectroscopic measurements. For convenience the transition and typical elements will be treated separately. IONISATION ENERGIES: TYPICAL ELEMENTS Changes down the group Table 2.1 gives data for Group I elements. The ionisation energies are all positive, i.e. energy is absorbed on ionisation. Several conclusions can be drawn from this table: 1. Energy must be supplied if these elements are to attain a noble gas configuration. 2. Loss of one electron gives the noble gas configuration; the very large difference between the first and second ionisation energies implies that an outer electronic configuration of a noble gas is indeed very stable. 3. Ionisation energy falls as the group is descended, i.e. as the size of the atom increases and hence the distance between the nucleus and the outer electron increases. 4. There is a marked contraction in size on the formation of an ion, the percentage contraction decreasing as the percentage loss in electrons decreases (for example Na -> Na 4 " involves loss of one of eleven electrons, Cs -> Cs + the loss of one of fifty-five electrons). Some values for Group II and III elements are shown in Tables 2.2 and 2.3 respectively.